Saltar al contenido principal
LibreTexts Español

2.3: Interacciones no covalentes

  • Page ID
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    Until now we have been focusing on understanding the covalent bonds that hold individual molecules together. We turn next to a brief review on the subject of non-covalent interactions between molecules, or between different functional groups within a single molecule.

    2.3A: Dipoles

    To understand the nature of noncovalent interactions, we first must return to covalent bonds and delve into the subject of dipoles. Many of the covalent bonds that we saw in section 1.5 – between two carbons, for example, or between a carbon and a hydrogen –involve the approximately equal sharing of electrons between the two atoms in the bond. This is because, in these examples, the two atoms have approximately the same electronegativity. Recall from general chemistry that electronegativity refers to "the power of an atom in a molecule to attract electrons to itself" (this is the definition offered by Linus Pauling, the eminent 20th-century American chemist who was primarily responsible for developing many of the bonding concepts that we have been learning).

    Quite often, however, we deal in organic chemistry with covalent bonds between two atoms with very different negativities, and in these cases the sharing of electrons is not equal: the more electronegative atom pulls the two electrons closer to itself. In the carbon-oxygen bond of an alcohol, for example, the two electrons in the s bond are closer to the oxygen than they are to the carbon, because oxygen is significantly more electronegative than carbon. The same is true for the oxygen-hydrogen bond, as hydrogen is slightly less electronegative than carbon, and much less electronegative than oxygen. 


    The result of this unequal sharing is what we call a bond dipole, or a polar covalent bond. A polar covalent bond has both negative and positive ends, or poles, where electron density is lower (the positive pole) and higher (the negative pole). The difference in electron density can be expressed using the Greek letter ∂ (delta) to express ‘partial positive’ and ‘partial negative’ charge on the atoms. ‘Dipole arrows’, with a positive sign on the tail, are also used to indicated the direction of higher electron density.

    The degree of polarity depends on the difference in electronegativity between the two atoms in the bond. Electronegativity is a periodic trend: it increases going from left to right across a row of the periodic table of the elements, and also increases as we move up a column. Therefore, oxygen is more electronegative than nitrogen, which is in turn more electronegative than carbon. Oxygen is also more electronegative than sulfur. Fluorine, in the top right corner of the periodic table, is the most electronegative of the elements, a property which we will see in later chapters can be exploited by researchers who are trying to work out the details of how organic reactions occur.

    Exercise 2.9

    Using the ideas that you have learned about atomic orbitals, rationalize the periodic trends in electronegativity.


    Most molecules contain both polar and nonpolar covalent bonds. Depending on the location of polar bonds and bonding geometry, molecules may posses an overall dipole moment. Water, as you probably recall, has a dipole moment that results from the combined dipoles of its two oxygen-hydrogen bonds. Fluoromethane also has an overall dipole moment. 


    Carbon dioxide, however, has polar bonds that pull in opposite directions, meaning that there is no overall molecular dipole moment.

    Exercise 2.10

    Which of the molecules below have overall dipole moments?



    2.3B: Ion-ion, dipole-dipole and ion-dipole interactions

    Arguably the simplest kind of non-covalent interaction is one that we have already discussed: the ionic bond formed by the electrostatic attraction between positively and  negatively charged species. One of the most common types of ion-ion interaction in biological chemistry is that between a magnesium ion (Mg+2) and a carboxylate or phosphate group.  2-phosphoglycerate, an intermediate in the breakdown of glucose, provides a good example in the way that it interacts with two Mg+2 ions while binding to an enzyme called enolase.


    Polar molecules – those with an overall dipole moment, such as acetone – will align themselves in such a way as to allow their respective positive and negative poles to interact with each other.  This is called a  ‘dipole-dipole interaction


    When a charged species (an ion) interacts favorably with a polar molecule or functional group, the result is called an ion-dipole interaction. A common example of ion-dipole interaction in biological organic chemistry is that between a metal cation, most often Mg+2 or Zn+2, and the partially negative oxygen of a carbonyl.


    Because the metal cation is very electronegative, this interaction has the effect of pulling electron density in the carbonyl double bond even further toward the oxygen side, increasing the partial positive charge on carbon.  As we shall later, this has important implications in terms of the reactivity of carbonyl groups.

    2.3C: van der Waals forces

    Nonpolar molecules such as hydrocarbons also are subject to relatively weak but still significant attractive intermolecular forces. Van der Waals forces (also called London dispersion forces, hydrophobic interactions, or nonpolar interactions) result from the constantly shifting electron density in any molecule. Even a nonpolar molecule will, at any given moment, have a weak, short-lived dipole.  This transient dipole will induce a neighboring nonpolar molecule to develop a corresponding transient dipole of its own, with the end result that a weak dipole-dipole force is formed. These van der Waals forces are relatively weak by nature, but are constantly forming and dissipating among closely-packed nonpolar molecules, and when added up the cumulative effect can become significant.  It is through van der Waals interactions that the nonpolar hydrocarbon regions of fat molecules hold together to form the lipid bilayer in cells and subcellular organelles (we will discuss this topic in more detail in section 2.4B)

    2.3D: Hydrogen bonds

    Hydrogen bonds are an extremely important form of noncovalent interaction. Hydrogen bonds result from the interaction between a hydrogen that is bound to a electronegative heteroatom – specifically a nitrogen or oxygen, or sometimes a fluorine – and a lone pair on a second nitrogen or oxygen.  Because a hydrogen is just a single proton with no inner shell (non-valence) electrons,  when it loses electron density in a polar bond it essentially becomes an approximation of a ‘naked’ proton, capable of forming a strong interaction with a lone pair on a neighboring oxygen or nitrogen atom.


    Hydrogen bonds are usually depicted with dotted lines in chemical structures. A group that provides a proton to a hydrogen bond is said to be acting as a hydrogen bond donor.  A group that provides an oxygen or nitrogen lone pair is said to be acting as a hydrogen bond acceptor. Many common organic functional groups can participate in the formation of hydrogen bonds, either as donors, acceptor, or both. A water molecule, for example, can be both a hydrogen bond donor and acceptor.  A carbonyl, as it lacks a hydrogen bound to an oxygen or nitrogen,  can only act as a hydrogen bond acceptor. 

    Exercise 2.11

    Draw figures that show the hydrogen bonds described below.

    1. A hydrogen bond between  methanol (donor) and water (acceptor). 
    2. A hydrogen bond between  methanol (acceptor) and water (donor).
    3. Two possible hydrogen bonds between methyl acetate and methylamine.


    Hydrogen bonds are the strongest type of noncovalent interaction, quite a bit stronger than a dipole-dipole interaction, but also much weaker than a covalent bond. The strength of hydrogen bonds has enormous implications in biology.  Genetic inheritance, for example, is based on a very specific hydrogen bonding arrangement between DNA bases: A ‘base-pairs’ with T, while G ‘base-pairs’ with C.

    image104.png       image106.png

    This page titled 2.3: Interacciones no covalentes is shared under a not declared license and was authored, remixed, and/or curated by Tim Soderberg.